Transcript Chapter 4 Atomic Structure

Chapter 4
Atomic Structure
• People have been thinking about the nature of
matter for a long time. The ancient Greeks
thought about matter and it wasn’t until the
late 19th century that an accepted theory was
arrived at.
• John Dalton in 1808 is credited with the
modern atomic theory. See page 103.
Fundamental Particles
Three fundamental particles make up atoms. The
following table lists these particles together with their
masses and their charges.
Particle
Mass (amu) Charge
Electron (e-)
0.00054858
-1
Proton (p,p+)
1.0073
+1
Neutron(n,n0)
1.0087
0
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The Discovery of Electrons
• Humphrey Davy in the early 1800’s passed
electricity through compounds and noted and
concluded that:
– the compounds decomposed into elements.
– compounds are held together by electrical forces.
• Michael Faraday in 1832-1833 realized that
the amount of reaction that occurs during
electrolysis is proportional to the electrical
current passed through the compounds.
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The Discovery of Electrons
• Cathode Ray Tubes experiments performed in the
late 1800’s & early 1900’s.
– Consist of two electrodes sealed in a glass tube containing
a gas at very low pressure.
– When a voltage is applied to the cathodes a glow discharge
is emitted.
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The Discovery of Electrons
• These “rays” are emitted from cathode (- end)
and travel to anode (+ end).
– Cathode Rays must be negatively charged!
• J.J. Thomson modified the cathode ray tube
experiments in 1897 by adding two adjustable
voltage electrodes.
– Studied the amount that the cathode ray beam
was deflected by additional electric field.
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The Discovery of Electrons
• Thomson used his modification to measure the
charge to mass ratio of electrons.
Charge to mass ratio
e/m = -1.75882 x 108 coulomb/g
• Thomson named the cathode rays electrons.
• Thomson is considered to be the “discoverer of
electrons”.
• TV sets and computer screens are cathode ray
tubes.
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Canal Rays and Protons
• Eugene Goldstein noted streams of positively charged particles in
cathode rays in 1886.
– Particles move in opposite direction of cathode rays.
– Called “Canal Rays” because they passed through holes (channels or
canals) drilled through the negative electrode.
• Canal rays must be positive.
– Goldstein postulated the existence of a positive fundamental
particle called the “proton”.
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Rutherford and the Nuclear Atom
• Ernest Rutherford directed Hans Geiger and
Ernst Marsden’s experiment in 1910.
– - particle scattering from thin Au foils
– Gave us the basic picture of the atom’s structure.
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Rutherford and the Nuclear Atom
• In 1912 Rutherford decoded the -particle
scattering information.
– Explanation involved a nuclear atom with electrons
surrounding the nucleus .
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Rutherford and the Nuclear Atom
Rutherford’s major conclusions from the -particle
scattering experiment
1. The atom is mostly empty space.
2. It contains a very small, dense center called the
nucleus.
3. Nearly all of the atom’s mass is in the nucleus.
4. The nuclear diameter is 1/10,000 to 1/100,000 times
less than atom’s radius.
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Rutherford and the Nuclear Atom
• Because the atom’s mass is contained in such
a small volume:
– The nuclear density is ~1015g/mL.
– This is equivalent to ~3.72 x 109 tons/in3.
– Density inside the nucleus is almost the same as a
neutron star’s density.
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Atomic Number
• The atomic number is equal to the number of
protons in the nucleus.
– Sometimes given the symbol Z.
– On the periodic table Z is the uppermost number in each
element’s box.
• In 1913 H.G.J. Moseley realized that the atomic
number determines the element .
– The elements differ from each other by the number of
protons in the nucleus.
– The number of electrons in a neutral atom is also equal to
the atomic number.
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Neutrons
• James Chadwick in 1932 analyzed the results
of -particle scattering on thin Be films.
• Chadwick recognized existence of massive
neutral particles which he called neutrons.
– Chadwick discovered the neutron.
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Mass Number and Isotopes
• Mass number is given the symbol A.
• A is the sum of the number of protons and neutrons.
– Z = proton number N = neutron number
– A=Z+N
• A common symbolism used to show mass and proton
numbers is
A
12
48
Z
6
20
 Can be shortened to this symbolism.
E for example
14
63
N, Cu,
C, Ca,
107
Ag, etc.
197
79
Au
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Mass Number and Isotopes
• Isotopes are atoms of the same element but with
different neutron numbers.
– Isotopes have different masses and A values but are the
same element.
• One example of an isotopic series is the hydrogen
isotopes.
1H
or protium is the most common hydrogen isotope.
• one proton and no neutrons
2H
or deuterium is the second most abundant hydrogen
isotope.
• one proton and one neutron
3H
or tritium is a radioactive hydrogen isotope.
• one proton and two neutrons
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Mass Number and Isotopes
• The stable oxygen isotopes provide another
example.
• 16O is the most abundant stable O isotope.
• How many protons and neutrons are in 16O?
8 protons and 8 neutrons
 17O
is the least abundant stable O isotope.
 How many protons and neutrons are in 17O?
8 protons and 9 neutrons
 18O
is the second most abundant stable O isotope.
How many protons and neutrons in 18O?
8 protons and 10 neutrons
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Mass Number and Isotopes
• Mass number is given the symbol A.
• A is the sum of the number of protons and neutrons.
– Z = proton number N = neutron number
– A=Z+N
• A common symbolism used to show mass and proton
numbers is
A
12
48
Z
6
20
 Can be shortened to this symbolism.
E for example
14
63
N, Cu,
C, Ca,
107
Ag, etc.
197
79
Au
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Mass Number and Isotopes
• Isotopes are atoms of the same element but with
different neutron numbers.
– Isotopes have different masses and A values but are the
same element.
• One example of an isotopic series is the hydrogen
isotopes.
1H
or protium is the most common hydrogen isotope.
• one proton and no neutrons
2H
or deuterium is the second most abundant hydrogen
isotope.
• one proton and one neutron
3H
or tritium is a radioactive hydrogen isotope.
• one proton and two neutrons
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Mass Number and Isotopes
• The stable oxygen isotopes provide another
example.
• 16O is the most abundant stable O isotope.
• How many protons and neutrons are in 16O?
8 protons and 8 neutrons
 17O
is the least abundant stable O isotope.
 How many protons and neutrons are in 17O?
8 protons and 9 neutrons
 18O
is the second most abundant stable O isotope.
How many protons and neutrons in 18O?
8 protons and 10 neutrons
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The Atomic Weight Scale and Atomic
Weights
• The atomic weight of an element is the
weighted average of the masses of its stable
isotopes
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Atomic Weight Scale and Atomic
Weights
Example 4-2: Naturally occurring Cu consists of 2
isotopes. It is 69.1% 63Cu with a mass of 62.9 amu, and
30.9% 65Cu, which has a mass of 64.9 amu. Calculate
the atomic weight of Cu to one decimal place.
atomic weight  (0.691)(62 .9 amu)  (0.309)(64 .9 amu)


 


63
Cu isotope
65
Cu isotope
atomic weight  63.5 amu for copper
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The Atomic Weight Scale and Atomic
Weights
Example 4-3: Naturally occurring chromium consists of
four isotopes. It is 4.31% 2450Cr, mass = 49.946 amu,
83.76% 2452Cr, mass = 51.941 amu, 9.55% 2453Cr, mass =
52.941 amu, and 2.38% 2454Cr, mass = 53.939 amu.
Calculate the atomic weight of chromium.
You do it!
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